Calculations Involving Masses Practice Questions

Question 1

3 marks

Use the relative atomic masses in the table to answer the questions below.

Element	H	C	O	Al	S	K
Relative atomic mass	1	12	16	27	32	39

(a)

Calculate the relative formula mass of potassium carbonate, $\mathrm{K_2CO_3}$ .

relative formula mass = ..............................................................

/ 1

(b)

Calculate the relative formula mass of aluminium sulfate, $\mathrm{Al_2(SO_4)_3}$ .

relative formula mass = ..............................................................

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Question 2

3 marks

Ammonium nitrate, $\mathrm{NH_4NO_3}$ , is used as a fertiliser because it supplies nitrogen to plants.

(relative atomic masses: H = 1, N = 14, O = 16)

(a)

Calculate the relative formula mass of ammonium nitrate, $\mathrm{NH_4NO_3}$ .

relative formula mass = ..............................................................

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(b)

Calculate the percentage by mass of nitrogen in ammonium nitrate.

percentage of nitrogen = .............................................................. %

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Question 3

4 marks

A gardener makes up a solution of potassium nitrate to feed plants.

She dissolves 8.0 g of potassium nitrate in water and makes the solution up to $400\ \mathrm{cm^3}$ .

(a)

Calculate the concentration of the solution in $\mathrm{g\ dm^{-3}}$ .

concentration = .............................................................. $\mathrm{g\ dm^{-3}}$

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(b)

The gardener needs $250\ \mathrm{cm^3}$ of potassium nitrate solution of the same concentration.

Calculate the mass of potassium nitrate she should dissolve to make this solution.

mass = .............................................................. g

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Question 4

3 marks

Chemists use the mole to count particles such as atoms, molecules and ions.

(relative atomic masses: C = 12, O = 16)

(a)

One mole of a substance contains the Avogadro constant number of particles.

Which of these is the value of the Avogadro constant?

A $6.02 \times 10^{-23}$
B $6.02 \times 10^{22}$
C $6.02 \times 10^{23}$
D $6.02 \times 10^{24}$

/ 1

(b)

Calculate the number of moles of carbon dioxide molecules in 11 g of carbon dioxide, $\mathrm{CO_2}$ .

number of moles = .............................................................. mol

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Question 5

3 marks

A student mixes lead(II) nitrate solution with potassium iodide solution in a conical flask fitted with a stopper. A yellow precipitate of lead(II) iodide forms.

The flask and its contents are weighed before the solutions are mixed.

(a)

State the reading on the balance after the reaction is complete.

reading = .............................................................. g

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(b)

Explain your answer to part (a).

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Question 0

4 marks

Tin reacts with oxygen when it is heated in air.

A student heated 2.38 g of tin in an open crucible until all the tin had reacted. The mass of tin oxide formed was 3.02 g.

(relative atomic masses: O = 16, Sn = 119)

(a)

Calculate the mass of oxygen that reacted with the tin.

mass of oxygen = .............................................................. g

/ 1

(b)

Calculate the empirical formula of this tin oxide.

You must show your working.

empirical formula = ..............................................................

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Question 0

5 marks

A hydrocarbon X contains 85.7% carbon and 14.3% hydrogen by mass.

(relative atomic masses: H = 1, C = 12)

(a)

Calculate the empirical formula of X.

You must show your working.

empirical formula = ..............................................................

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(b)

The relative molecular mass of X is 56.

Deduce the molecular formula of X.

molecular formula = ..............................................................

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Question 0

4 marks

The empirical formula of a compound shows the simplest whole number ratio of the atoms of each element in the compound.

(relative atomic masses: O = 16, P = 31)

(a)

Give the empirical formula of glucose, $\mathrm{C_6H_{12}O_6}$ .

empirical formula = ..............................................................

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(b)

Give the empirical formula of butane, $\mathrm{C_4H_{10}}$ .

empirical formula = ..............................................................

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(c)

An oxide of phosphorus has the empirical formula $\mathrm{P_2O_5}$ and a relative molecular mass of 284.

Deduce the molecular formula of this oxide.

molecular formula = ..............................................................

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Question 0

4 marks

Calcium oxide (quicklime) is made in a lime kiln by heating calcium carbonate.

$\mathrm{CaCO_3} \rightarrow \mathrm{CaO} + \mathrm{CO_2}$

(relative atomic masses: C = 12, O = 16, Ca = 40)

(a)

Calculate the maximum mass of calcium oxide that can be made from 750 kg of calcium carbonate.

mass of calcium oxide = .............................................................. kg

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(b)

Calculate the mass of carbon dioxide released when 750 kg of calcium carbonate is completely decomposed.

mass of carbon dioxide = .............................................................. kg

/ 1

Question 0

4 marks

Mass changes can appear to happen when reactions are carried out in open containers.

(a)

A student heats a piece of copper in an open crucible. The mass of the solid in the crucible increases from 1.27 g to 1.59 g.

Explain why the mass increases.

/ 2

(b)

The student then adds marble chips to dilute hydrochloric acid in an open conical flask standing on a balance. The reading on the balance decreases during the reaction.

Explain why the reading decreases, and state what would happen to the reading if the flask were sealed with a stopper.

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Question 0

5 marks

(relative atomic masses: O = 16, Al = 27)

(Avogadro constant = $6.02 \times 10^{23}\ \mathrm{mol^{-1}}$ )

(a)

Calculate the number of moles of aluminium atoms in 2.7 g of aluminium.

number of moles = .............................................................. mol

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(b)

Calculate the number of aluminium atoms in 2.7 g of aluminium.

Give your answer in standard form.

number of atoms = ..............................................................

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(c)

Calculate the mass of $1.505 \times 10^{23}$ molecules of oxygen, $\mathrm{O_2}$ .

mass = .............................................................. g

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Question 0

4 marks

A student evaporates $25.0\ \mathrm{cm^3}$ of a sodium chloride solution to dryness in an evaporating basin. The mass of solid sodium chloride left behind is 0.85 g.

(a)

Calculate the concentration of the sodium chloride solution in $\mathrm{g\ dm^{-3}}$ .

concentration = .............................................................. $\mathrm{g\ dm^{-3}}$

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(b)

Calculate the volume, in $\mathrm{cm^3}$ , of the same solution that contains 5.1 g of sodium chloride.

volume = .............................................................. $\mathrm{cm^3}$

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Question 0

5 marks

Iron reacts with sulfur when the mixture is heated.

$\mathrm{Fe} + \mathrm{S} \rightarrow \mathrm{FeS}$

A student heats 14.0 g of iron filings with 4.0 g of sulfur until no further reaction occurs.

(relative atomic masses: S = 32, Fe = 56)

(a)

Show that sulfur is the limiting reactant in this experiment.

/ 2

(b)

Calculate the maximum mass of iron(II) sulfide, FeS, that can be formed.

mass of FeS = .............................................................. g

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(c)

Explain why the mass of iron(II) sulfide formed is controlled by the mass of sulfur used.

/ 1

Question 0

6 marks

A student uses the apparatus shown to find the empirical formula of magnesium oxide.

(a)

Describe how the student should carry out this experiment to obtain accurate results, including the measurements that must be made.

/ 4

(b)

The student's results gave an empirical formula of $\mathrm{Mg_5O_4}$ instead of $\mathrm{MgO}$ .

Suggest two reasons why the ratio of oxygen to magnesium in the student's result was too low.

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Question 0

4 marks

A student reacted magnesium with dilute hydrochloric acid and measured the masses of all the substances involved.

Substance	Mass / g
magnesium used	2.4
hydrogen chloride reacted	7.3
magnesium chloride formed	9.5
hydrogen formed	0.2

(relative atomic masses: H = 1, Mg = 24, Cl = 35.5)

(a)

Calculate the number of moles of each substance in the table.

magnesium = .............................................................. mol

hydrogen chloride = .............................................................. mol

magnesium chloride = .............................................................. mol

hydrogen = .............................................................. mol

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(b)

Use your answers to part (a) to deduce the balanced equation for this reaction.

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Question 0

5 marks

Hydrated copper(II) sulfate has the formula $\mathrm{CuSO_4}\cdot x\mathrm{H_2O}$ .

A student heated 5.00 g of hydrated copper(II) sulfate in an evaporating basin until all the water of crystallisation had been driven off. The mass of anhydrous copper(II) sulfate left was 3.20 g.

(relative atomic masses: H = 1, O = 16, S = 32, Cu = 63.5)

(a)

Calculate the mass of water driven off.

mass of water = .............................................................. g

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(b)

Calculate the value of $x$ in $\mathrm{CuSO_4}\cdot x\mathrm{H_2O}$ .

You must show your working.

$x$ = ..............................................................

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(c)

Describe how the student could be sure that all of the water of crystallisation had been driven off.

/ 1

Question 0

5 marks

Methane burns completely in oxygen.

$\mathrm{CH_4} + 2\mathrm{O_2} \rightarrow \mathrm{CO_2} + 2\mathrm{H_2O}$

(relative atomic masses: H = 1, C = 12, O = 16)

(Avogadro constant = $6.02 \times 10^{23}\ \mathrm{mol^{-1}}$ )

(a)

Calculate the mass of oxygen needed to burn 4.0 g of methane completely.

mass of oxygen = .............................................................. g

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(b)

Calculate the number of carbon dioxide molecules produced when 4.0 g of methane burns completely.

Give your answer in standard form.

number of molecules = ..............................................................

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Question 0

5 marks

Urea, $\mathrm{CO(NH_2)_2}$ , and ammonium sulfate, $\mathrm{(NH_4)_2SO_4}$ , are both used as nitrogen fertilisers.

(relative atomic masses: H = 1, C = 12, N = 14, O = 16, S = 32)

(a)

Show that the percentage by mass of nitrogen in urea, $\mathrm{CO(NH_2)_2}$ , is 46.7% to 3 significant figures.

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(b)

Calculate the percentage by mass of nitrogen in ammonium sulfate, $\mathrm{(NH_4)_2SO_4}$ .

Give your answer to 3 significant figures.

percentage of nitrogen = .............................................................. %

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(c)

Calculate the mass of nitrogen, in kg, contained in a 25 kg bag of urea.

mass of nitrogen = .............................................................. kg

/ 1

Atomic Structure

States of Matter & Mixtures

Chemical changes

Extracting metals & equilibria

Separate Chemistry 1

Groups in the periodic table

Rates of reaction & energy changes

Fuels & Earth science

Separate Chemistry 2

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